THE ATOM

1. THE ATOM AP CHEMISTRY 1. Take out a pencil and sit in testing seats

2. Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

3. Dalton’s Postulates Each element is composed of extremely small particles called atoms.

4. Dalton’s Postulates All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

5. Dalton’s Postulates Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

6. Dalton’s Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

7. Law of Constant CompositionJoseph Proust (1754–1826) • Also known as the law of definite proportions. • The elemental composition of a pure substance never varies.

8. Multiple Proportions • The law states that when chemical elements combine, they do so in a ratio of small whole numbers (Ex. Carbon and oxygen react to form carbon monoxide (CO) or carbon dioxide (CO2) but not CO13 • If two elements form more than one compound between them, the ratio of the masses of the second element to a mass of the first elements will also be in small whole numbers

9. Law of Conservation of Mass The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

10. The Discovery of an Electron • Early scientists knew about charges and in fact, Benjamin Franklin gave the names positive and negative to the two different charges. • But Michael Faraday (1791-1867) discovered electrons as “cathode rays” by applying a high voltage to the ends of a cathode ray tube. • The electrons emitted provide an image of their path when they strike a fluorescence zinc sulfide screen.

11. The Discovery of an Electron • J.J. Thomson (1856-1940) used a specially designed cathode ray tube to apply both electric and magnetic fields simultaneously to the beam of cathode rays. • By balancing the effect of the electric field against that of the magnetic field he was able to calculate the charge to mass (e/m) ratio for the particles in the beam. • Thomson’s experiments also demonstrated that electrons had a negative charge. In addition, he obtained the same mass to charge ratio with twenty different metals. • These results suggested that electrons are present in atoms of all elements.

12. The Atom, circa 1900: • “Plum pudding” model, put forward by Thompson. • Positive sphere of matter with negative electrons imbedded in it.

13. The Millikan Oil Drop Experiment • American Physicist Robert Millikan (1868 - 1953) performed an experiment in which he sprayed oil droplets into a chamber from an atomizer. The oil droplets were allowed to settle slowly towards the bottom of the chamber. • Millikan had two charged plates on the top and bottom of the chamber which allowed him to control the rate at which the oil droplets fell. • By varying the voltage on the plates, he was able to just stop the oil droplets from falling.

14. As a result of his experiment, Millikan was able to calculate the following charges on the oil droplets • What can you conclude about the charges which Millikan obtained? • They are all integer multiples of -1.6 x 10-19 C. Since the charge on an electron is the smallest possible charge, it is called an “elementary charge.”

15. RADIOACTIVITY • The spontaneous emission of radiation by an atom. • First observed by Henri Becquerel. • Also studied by Marie and Pierre Curie.

16. RADIOACTIVITY • Three types of radiation were discovered by Ernest Rutherford: •  particles •  particles •  rays

17. Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

18. The Nuclear Atom Since some particles were deflected at large angles, Thompson’s model could not be correct.

19. The Nuclear Atom • Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. • Most of the volume of the atom is empty space.

20. The Rutherford Scattering Experiment • In 1910, Ernest Rutherford (1871 - 1937) and an associate, Hans Geiger (1882 - 1945) and a student, Ernest Marsden (1889 - 1970) performed an experiment in which they bombarded various metal foils with alpha particles. • They observed that most of the alpha particles passed through the metal foil with little or no deflection. Some of the alpha particles actually bounced back towards the alpha source.

21. An Enlarged View of Rutherford’s Experiment

22. CONCLUSIONS As a result of Rutherford’s experiment, scientists were able to conclude two important features of the atom: 1. The atom is mostly empty space. 2. The atom’s positive charge is concentrated in a central core within the atom.

23. Some Properties of the Gold Atom What can you conclude about the relative size of the gold atom, nucleus, and an electron? If an atom were the size of a football field (- 100 meters), the nucleus would be about 3 millimeters in diameter (the size of a BB) and an electron would be roughly half the size of the nucleus. That is, the atom is mostly empty space.

24. Some Properties of the Gold Atom • In other words, when Rutherford fired alpha particles at the gold foil, it would be like trying to hit a BB in the middle of a football field with a BB gun from many miles away, nearly impossible. • Rutherford also observed that only one in twenty thousand alpha particles were significantly deflected as they passed through the gold foil. Explain Rutherford’s observation in terms of the relative size of the atom and the nucleus. • What comparison can you make about the mass of the nucleus compared to the mass of the atom? • Almost all of the mass of an atom is contained in the nucleus.

25. ATOMIC STRUCTURE • The atom consists of two parts: • The positively charged nucleus containing most the atom’s mass. • Negatively charged electrons found in the empty space around the nucleus.

26. Other Subatomic Particles • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932.

27. CONCLUSION QUESTIONS Give It Some Thought Questions on pg. 41, 44 Go Figure Questions on pg. 42, 43

28. THE ATOM AP CHEMISTRY DO NOW:1. If an atom has 15 protons, how many electrons does it have?2. Where do the protons reside in an atom?3. The diameter of a US dime is 17.9 mm, and the diameter of a silver atom is 2.88Å. How many silver atoms could be arranged side by side across the diameter of a dime?Equivalence statements: 1Ag atom = 2.88Å1Å = 1 x10 -10 m

29. SOME IMPORTANT ATOMIC STRUCTURE TERMS Atomic number - The atomic number equals the number of protons in the nucleus. The atomic number determines the element. Nucleons - particles in the nucleus (protons plus neutrons) Nuclear charge - The charge on the nucleus (the nuclear charge equals the number of protons) Mass number - The mass number (an integer) is the sum of the neutrons plus protons. The mass number also equals the number of nucleons. Isotopes - Isotopes are atoms of the same element (same atomic number) with a different number of neutrons. Nuclide - A particular type of atom having a characteristic nucleus. Ion - An ion is a charged particle.

30. Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it.

31. IMPORTANT SUBATOMIC PARTICLES

32. Give the number of subatomic particles in each of the following atoms: This particle is called carbon - 14. Protons: _______ Nuclear charge: ______ Neutrons: _______ Electrons: _______ Number of nucleons: ______

33. Give the number of subatomic particles in each of the following atoms: Protons: _______ Neutrons: _______ Electrons: _______ Mass number: ______ What is the name of this particle? ______________

34. A certain particle has 11 protons, 13 neutrons, and 10 electrons. 1. What is the name of the element? 2. What is the mass number of this particle? 3. How many nucleons are found in this particle? 4. What is the overall charge on this particle? 5. Is this particle an atom or an ion? 6. Give the symbol for the nucleus of this particle.

35. Answer each of the following questions based on the following symbol: 1. What is the mass number of this particle? 2. How many protons, neutrons, and electrons are found in this particle? 3. What is the nuclear charge on this element?

36. Answer each of the following questions based on the following symbol: 1. What is the mass number of this particle? 2. How many protons, neutrons, and electrons are found in this particle? 3. What is the nuclear charge on this element?

37. Isotopes of Hydrogen Give the number of protons, neutrons, and electrons for each of the isotopes of hydrogen listed above.

38. Which of the following nuclides are isotopes?

39. Properties of Heavy Water and Ordinary Water

40. Units of Atomic Mass An atomic mass unit (amu or u) is defined as exactly 1/12 the mass of a carbon-12 atom. or the mass of a carbon-12 atom is exactly 12.000 . . . amu. The atomic masses of all the other elements are determined by comparing their masses with carbon-12.

41. Units of Atomic Mass For example, the mass of an average hydrogen atom is 8.400% the mass of a carbon-12 atom. Thus, the mass of an average hydrogen atom is 0.08400 x 12 amu = 1.008 amu. Likewise, an average atom of magnesium has a mass which is 2.0254 times the mass of a carbon- 12 atom. Therefore, the mass of an average magnesium atom is 2.0254 x 12 amu = 24.305 amu.

42. COMPARISON OF ATOMIC MASS SCALES

43. Determining the Mass of Atoms The mass of atoms is determined by use of a mass spectrograph. The sample of matter is ionized in a vacuum chamber. The resulting positive ions are then accelerated by means of a negatively charged screen. Most of the ions pass through the screen, though a slit to focus the ion beam, and then into a magnetic field. By varying the accelerating voltage, the ions can be made to strike the detector. Velocities up to 150,000 miles/sec can be obtained with a voltage of 400-4000 volts.

44. Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

45. Atomic Mass • Most accurate means for determining atomic weights • A gaseous sample is introduced and bombarded by a stream of high –energy electrons • Collisions between the electrons and the atoms of molecules of the gas produce positively charged particles that are then accelerated toward a negatively charged grid .

46. Atomic Mass • After particles pass through the grid, they encounter two slits that allow only a narrow beam of particles to pass • This beam then passes between the poles of a magnet, which deflects the particles into a curved path • Particles with the same charge, the extent of deflection depends on mass – the more massive the particle, the less the deflection

47. Atomic Mass • Particles are separated by according to their masses • By hanging the strength of the magentic field or the accelerating voltage on the grid, charged particles of various masses can be selected to enter the detector • A graph of the intensity of the detector signal versus particle atomic mass is called a mass spectrum

48. Atomic Mass • Analysis of a mass spectrum gives both the masses of the charged particles reaching the detector and their relative abundances, which are obtained from signal intensities • Knowing atomic mass and abundance of each isotope allows us to calculate the atomic weight of an element

49. Atomic Mass • Used extensively today to identify chemical compounds and analyze mixtures of substances • Any molecules that loses electrons can fall apart, forming an array of positively charged fragments • Mass spec measures the masses of these fragments, producing a chemical “fingerprint” of the moelcule and providing clues about how the atoms were connected in the original molecule • Application: used to determine the molecular structure of a newly synthesized compound or to identify a pollutant in the environment